Electrochemical Stability of LiMF6 (M = P, As, Sb) in
Tetrahydrofuran and Sulfolane
C. Nanjundiah,* J. L. Goldman,* L. A. Dominey,* and V. R. Koch*
Covalent Associates, Incorporated, Woburn, Massachusetts 01801
ABSTRACT
LiAsF6 is the current state-of-the-art lithium salt for use in m a n y p r o t o t y p e r e c h a r g e a b l e systems. T h e instability
and potential t o x i c i t y of LiAsF6 has led to the search for n e w l i t h i u m salts. To this end, w e investigated t h e e l e c t r o c h e m i cal stability of LiSbF6 and LiPF~ in aprotic organic solvents. E l e c t r o c h e m i c a l Studies, c o n d u c t i v i t y m e a s u r e m e n t s , and
open-circuit stability tests w e r e c o n d u c t e d on LiSbF6 in t e t r a h y d r o f u r a n and sulfolane. Cyclic v o l t a m m e t r i c studies of
SbF6 w e r e c o m p a r e d w i t h those of AsF6- and PF6- anions. Sb(V) was r e d u c e d to Sb(III), and t h e n to Sb(0). A s F c was red u c e d to AsF3, w h i l e no r e d u c t i o n of PF6- was observed. R e d u c t i o n p r o d u c t s of AsF6 and S b F C passivated the glassy carb o n electrode, p r e s u m a b l y due to L i F precipitation. P e a k potentials w e r e o b s e r v e d to shift in the positive direction as a
f u n c t i o n of concentration. This was a c c o u n t e d for on the basis of a follow-up c h e m i c a l reaction.
B e c a u s e of its c o m m e r c i a l availability in h i g h purity,
LiAsF~ is one of the m o s t popular s u p p o r t i n g electrolytes
u s e d t o d a y in a m b i e n t t e m p e r a t u r e r e c h a r g e a b l e Li batteries. Nevertheless, the AsF6- anion is k n o w n to be c h e m ically and e l e c t r o c h e m i c a l l y u n s t a b l e in a r e d u c i n g envir o n m e n t (1-3). Of greater potential c o n c e r n for t h e
c o n s u m e r m a r k e t is the t o x i c i t y of As(0), one of m a n y
A s F c r e d u c t i o n p r o d u c t s (4). In o r d e r to better u n d e r s t a n d
t h e c h e m i s t r y of AsF~- anions in particular, and M F C
anions in general, w e set out to e x p l o r e the e l e c t r o c h e m i s try of P F C , AsF6-, and SbF6 in c o m m o n aprotic organic
solvents. To this end, cyclic v o l t a m m e t r y (CV) and polaro g r a p h y w e r e c o n d u c t e d at glassy c a r b o n (GC) and the
d r o p p i n g m e r c u r y electrode (DME), respectively. In addition, the stability of the SbF6 c o n t a i n i n g electrolytes w i t h
Li at open-circuit voltage (OCV) at 70~ was studied.
Experimental
E l e c t r o c h e m i c a l e x p e r i m e n t s Were c o n d u c t e d at 25~
u n d e r an Ar a t m o s p h e r e in a V a c u u m - A t m o s p h e r e s Corp o r a t i o n dry b o x e q u i p p e d w i t h a M o d e l HE-493 dri-train~
Electrolytes w e r e p r e p a r e d from distilled or H P L C - g r a d e
solvents treated with Li-Hg a m a l g a m and filtered t h r o u g h
0.7~t W h a t m a n G F / F glass filters to r e m o v e the r e d u c e d impurities. The purity of t h e s e electrolytes was assessed by
CV and gas p h a s e c h r o m a t o g r a p h y .
S a m p l e s of electrolyte and freshly s c r a t c h e d Li foil w e r e
i n c u b a t e d in Teflon-lined screw-cap culture t u b e s (Corning, C9826) at 70 ~ -+ 2~ Visual observations of electrolyte
c o n d i t i o n s w e r e m a d e on a day-to-day basis.
Materials.--Lithium h e x a f l u o r o a n t i m o n a t e (LiSbF6)
was p r e p a r e d for us by the R e s e a r c h and P r o d u c t i v i t y
Council (RPC), N e w B r u n s w i c k , Canada, and was also obt a i n e d f r o m Ozark-Mahoning, Tulsa, Oklahoma. The R P C
material was of h i g h e r purity, b e i n g twice recrystallized
f r o m acetonitrile. It did not p o l y m e r i z e t e t r a h y d r o f u r a n
(THF). T h e Ozark-Mahoning material p o l y m e r i z e d T H F
w i t h i n 2 or 3h, suggesting significant SbF5 contamination.
L i t h i u m h e x a f l u o r o a r s e n a t e (LiAsF6) (La R o c h e Industries, I n c o r p o r a t e d , Decatur, Georgia), l i t h i u m hexafluor o p h o s p h a t e (LiPF6) (Morita K a g a k u C o m p a n y , Japan),
and t e t r a b u t y l a m m o n i u m f l u o r o b o r a t e (TBAFB) (Southw e s t Analytical Chemicals, I n c o r p o r a t e d , Austin, Texas)
w e r e all u s e d as received. T e t r a h y d r o f u r a n (>99.9%) was
o b t a i n e d f r o m F i s h e r Scientific. T w o - m e t h y l t e t r a h y d r o furan (2-MeTHF) and sulfolane (S), o b t a i n e d f r o m A l d r i c h
C h e m i c a l C o m p a n y , w e r e dried over Call2 and distilled
u n d e r N2 at r e d u c e d pressure. E t h y l e n e c a r b o n a t e (98%)
was o b t a i n e d f r o m A l d r i c h and purified by drying over
P205 and v a c u u m distillation. A n t i m o n y (III) fluoride
(SbF3) was 98% p u r e from A l d r i c h and u s e d as received. Li
foil (10 mil) was f r o m F o o t e Mineral, Exton, Pennsylvania.
A n t i m o n y (V) fluoride s u p p o r t e d on 50% g r a p h i t e was ob* Electrochemical Society Active Member.
t a i n e d from Alfa P r o d u c t s (Danvers, Massachusetts) and
was u s e d as received.
Electrochemical cell and instrumentation.--A t h r e e - n e c k
25 ml r o u n d b o t t o m flask was u s e d as an e l e c t r o c h e m i c a l
cell. T h e w o r k i n g electrodes w e r e a 3 m m d i a m glassy carbon (GC) disk (Tokai) sealed in P y r e x and a d r o p p i n g merc u r y electrode (DME). The GC disk was p o l i s h e d w i t h 0.3~t
a l u m i n a (Fisher) before and b e t w e e n scans. A typical
open-circuit potential drop t i m e of the D M E (Sargent
Welch) was 6s and the mass rate of m e r c u r y was 0.982
mg/s. T h e co'unterelectrode was fabricated f r o m Ni foil.
T e n mil Li sheet was u s e d as a reference e l e c t r o d e and was
d i p p e d directly into the b u l k electrolyte. All potentials are
q u o t e d w i t h respect to the Li/Li § couple.
CV data w e r e a c q u i r e d w i t h either a B A S CV-1B potentiostat/wave g e n e r a t o r c o u p l e d to an H P M o d e l 7035B X-Y
recorder, or a P A R C 273 potentiostat/galvanostat interfaced w i t h an IBM P C A T computer. S o f t w a r e written inh o u s e afforded data acquisition and i n s t r u m e n t control.
H a r d copies of CV traces w e r e o b t a i n e d via an H P M o d e l
7040A digital plotter.
C o n d u c t i v i t y m e a s u r e m e n t s w e r e o b t a i n e d at 1 kHz
w i t h a YSI 3400 dip-type c o n d u c t i v i t y cell on a G e n Rad
1650A i m p e d a n c e bridge. The cell c o n s t a n t was o b t a i n e d
by m e a s u r i n g t h e c o n d u c t i v i t y of standard KC1 solutions
and u s i n g the r e p o r t e d specific c o n d u c t i v i t i e s (5). The acc u r a c y of our m e a s u r e m e n t s was assessed by c o m p a r i s o n
w i t h an 0.87M L i A s F d p r o p y l e n e c a r b o n a t e electrolyte
w h o s e c o n d u c t i v i t y has b e e n p u b l i s h e d (6). The values
agree w i t h i n 10%.
Results and Discussion
Electrolyte reactivity toward Li at open circuit.--LiSbF6
reactivity t o w a r d Li m e t a l in T H F and sulfolane was investigated. A thick, black deposit was f o r m e d on Li in all
LiSbF6 c o n t a i n i n g electrolytes. The film t h i c k n e s s was obs e r v e d to g r o w as a f u n c t i o n of time. CV studies (vide
infra) indicated that SbF6- was r e d u c e d on GC prior to Li +
reduction. This indicates that Li foil s p o n t a n e o u s l y red u c e s S b F C . Most p r o b a b l y Sb(0) and L i F are d e p o s i t e d
on t h e Li surface.
Conductivity.--To m i n i m i z e o v e r p o t e n t i a l due to IR
d r o p in a battery, the use of a highly c o n d u c t i n g electrolyte
is desirable. To this end, w e m e a s u r e d t h e specific conductivity o f LiSbF6 as a f u n c t i o n of salt concentration. F i g u r e
1 c o m p a r e s the c o n d u c t i v i t y of LiSbF6 w i t h that of LiAsF6
in T H F and sulfolane. Identical conductivities, w i t h i n exp e r i m e n t a l error, w e r e o b t a i n e d f r o m two different
sources of LiSbF6 u s i n g two calibrated c o n d u c t i v i t y cells.
Cyclic voltammetric studies.--CV studies w e r e u s e d to
ascertain the e l e c t r o c h e m i c a l w i n d o w o f the electrolytes
c o n t a i n i n g t h e s e M F C anions in T H F and sulfolane. As
was n o t e d p r e v i o u s l y by Dey and R u d d (7), anodic limits
are critically i m p o r t a n t w i t h respect to setting t h e maxi2914
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Vol. 135, No. 12
LiMF6 IN TETRAHYDROFURAN
~8
Y
30
17
15
15
14
13
I
AND SULFOLANE
l
t
[
[
3.2
2,8
2915
L
t
I
I
J
l
i
I
20
i0
LiAsFs/THF/'/'Y']"
//
x
/
8
10
LiSbF6/THF ,~
6
2~
5
4
30
Li AsF6/ S Z~ZZZ=~
:z
~.-------~~
3
2
Ls
[
.....
I
4
I__
3,6
I
I
2J4
2
ii,6
[
1.2
01,8
01,4
~,
I
(VOLTS)
Fig. 2. CV curve of 1 mM LiPF6 in a 0.1M TBAFB/THF at GC. Scan
rate = 100 mV/s, area of GC = 0.07 cm2.
POTEtITIAL VS Li
0
i
0.4
i
1
0.8
i
J
1.2
coNcENTRATION(Ill
Fig. 1. Specific canductivities of LiSbF6 and LiAsF6 in THF and sulfolane at 25~
mum charging potentials in a battery. The glassy carbon
(GC) e l e c t r o d e w a s u s e d to i n v e s t i g a t e t h e a n o d i c limits.
T h e a n o d i c p o t e n t i a l a t a c u r r e n t d e n s i t y of 100 ~ A / c m 2
(Eal00) w a s u s e d to d e f i n e t h e a n o d i c limit, as h a s b e e n d o n e
p r e v i o u s l y (8). O n e c a n i n f e r f r o m T a b l e I t h a t t h e o x i d a t i o n l i m i t s of e l e c t r o l y t e s c o n t a i n i n g L i S b F 6 a n d LiPF6
c o m p a r e well w i t h t h a t of t h e m o s t c o m m o n l y u s e d
LiAsF~.
C a t h o d i c l i m i t s are also i m p o r t a n t s i n c e a n y species, particularly anions, reducing prior to Li + may be irreversibly
d e g r a d e d as t h e cell is d i s c h a r g e d . T h i s n e c e s s a r i l y dec r e a s e s t h e n u m b e r of soluble, c o n d u c t i n g a n i o n s p e c i e s
a v a i l a b l e for cycling. T h e r e f o r e , a n u n d e r s t a n d i n g o f a n i o n
r e d u c i b i l i t y is i m p o r t a n t . We e m p l o y e d a G C w o r k i n g elect r o d e to i n v e s t i g a t e t h e c a t h o d i c r e a c t i v i t y of t h e s e a n i o n s .
T h e c a t h o d i c l i m i t s at E~00 a r e also s h o w n i n T a b l e I.
LiPF6.--No r e d u c t i o n or o x i d a t i o n w a v e s w e r e o b s e r v e d
for LiPF6 i n 0.1M T B A F B / T H F e l e c t r o l y t e s w i t h i n t h e
s c a n n e d v o l t a g e l i m i t s (Fig. 2). H o w e v e r , it h a s b e e n rep o r t e d (9) t h a t PF6 is t h e r m a l l y u n s t a b l e a n d d i s p r o p o r t i o n a t e s to give PF~ a n d F-, t h e f o r m e r o f w h i c h i n i t i a t e s
t h e p o l y m e r i z a t i o n of m a n y e t h e r s o l v e n t s .
LiAsF6.--A w e l l - d e f i n e d r e d u c t i o n p e a k w a s o b s e r v e d
for AsF6- as s h o w n i n Fig. 3. T h i s figure also s h o w s t h a t t h e
p e a k p o t e n t i a l shifts p o s i t i v e as t h e c o n c e n t r a t i o n of
LiAsF6 is i n c r e a s e d . H o w e v e r , t h e p e a k c u r r e n t d o e s n o t
i n c r e a s e p r o p o r t i o n a l l y . B a s e d o n a p e a k c u r r e n t of 63 ~tA
at 9 m M of LiAsF6, o n e w o u l d e x p e c t a c u r r e n t o f 14.7 m A
at 2.1M. B u t t h e o b s e r v e d c u r r e n t w a s o n l y 138 #A. F u r ther, no peak was seen on successive cycling of the potential b e t w e e n 3 a n d 0.5V. I n fact, t h e p e a k c o m p l e t e l y disapp e a r s a f t e r t h e t h i r d cycle. T h e s e o b s e r v a t i o n s i n d i c a t e
t h a t t h e p r o d u c t of e l e c t r o c h e m i c a l r e d u c t i o n o f A s F C
p a s s i v a t e s t h e e l e c t r o d e s . A p o s s i b l e e l e c t r o c h e m i c a l red u c t i o n m e c h a n i s m c a n b e w r i t t e n as follows
A s F C + 2e- ~ AsF3 + 3 F -
[1]
F - + Li + ~ L i F
[2]
is f u r t h e r s u p p o r t e d b y t h e fact t h a t t h e a d d i t i o n o f s m a l l
a m o u n t s (150 to 450 p p m ) o f w a t e r i n c r e a s e s t h e p e a k
h e i g h t . T h i s i n c r e a s e o c c u r s b e c a u s e L i F is m o r e s o l u b l e i n
a q u e o u s T H F t h a n i n T H F itself. F o r m a t i o n of L i F w a s obs e r v e d in a n earlier s t u d y w h e n L i w a s c y c l e d in L i A s F d
T H F e l e c t r o l y t e (2). T h i s s t u d y also r e v e a l e d t h a t As(V)
w a s r e d u c e d to As (III) o n t h e l i t h i u m electrode.
I t is p e r t i n e n t to p o i n t o u t t h e effect of L i F f o r m a t i o n o n
t h e i-E c u r v e (Fig. 3). S i n c e t h e p e a k p o t e n t i a l (Ep) is t h e res u l t of a n e l e c t r o c h e m i c a l p r o c e s s w h i c h l e a d s to r a p i d
electrode passivation, one cannot ascribe any thermodyn a m i c s i g n i f i c a n c e to t h e Ep values. H e n c e , w e c h o s e a pot e n t i a l a t t h e r i s i n g p a r t o f t h e i-E c u r v e at a c o n v e n i e n t .
c u r r e n t d e n s i t y , i.e., 210 ~AJcm 2. T h i s p o t e n t i a l , a f t e r I R
c o r r e c t i o n (10), s h i f t e d i n t h e a n o d i c d i r e c t i o n w h e n [Li §
i n c r e a s e d . T h i s o b s e r v a t i o n is i n c o n s o n a n c e w i t h t h e wellk n o w n effect o f c h e m i c a l f o l l o w - u p r e a c t i o n s o n t h e proc e e d i n g e l e c t r o c h e m i c a l s t e p (an E C m e c h a n i s m ) . H o w ever, a q u a n t i t a t i v e c a l c u l a t i o n o f s h i f t i n p o t e n t i a l as a
f u n c t i o n o f Li § c o n c e n t r a t i o n c a n n o t b e m a d e d u e to t h e
l a c k of a c t i v i t y d a t a at t h e s e c o n c e n t r a t i o n s .
LiSbF6.--A C V o f 3.9 m M L i S b F 6 i n T H F is s h o w n i n Fig.
4. T h e v e r y first cycle is s h o w n as a solid l i n e a n d t h e seco n d cycle as a d o t t e d line. O n s c a n n i n g n e g a t i v e f r o m O C V
(3V), t w o r e d u c t i o n p e a k s at 1.1V ( p e a k 1) a n d 0.9V ( p e a k 2)
were observed. On reversing the scan direction, one
a n o d i c p e a k a t 2.8V ( p e a k 3) w a s o b s e r v e d , i n d i c a t i n g t h a t
t h i s o x i d a t i o n is t h e r e s u l t of p r o d u c t g e n e r a t e d f r o m
Sb(V) r e d u c t i o n o c c u r r i n g u n d e r p e a k s 1 a n d 2. S i n c e
p e a k s 1 a n d 2 overlap, it is n o t p o s s i b l e to a s c e r t a i n
w h e t h e r p e a k 1 or p e a k 2 is r e s p o n s i b l e for p e a k 3. O n cont i n u i n g t h e s c a n to t h e s e c o n d cycle, a n o t h e r r e d u c t i o n
p e a k ( p e a k 4) w a s s e e n at 2.2V ( d o t t e d curve). O n r e v e r s i n g
t h e s e c o n d s c a n at 1.5V, a w e l l - d e f i n e d s t r i p p i n g p e a k
~ 5~,A
S i n c e L i F is i n s o l u b l e i n T H F , it will p r e c i p i t a t e o n t h e
e l e c t r o d e , t h e r e b y p a s s i v a t i n g t h e e l e c t r o d e surface. T h i s
Table I. Electrochemical window of various electrolytes at GC
electrode at 25~
Salt
Solvent
Conc. (M)
~Eam0(V)
ECru0(V)
LiSbFe
THF
Sulfolane
THF
Sulfolane
THF
Sulfolane
1.0
0.8
1.0
0.8
0.001a
0.001a
4.10
4.50
4.25
4.69
4.4
4.8
1.80
1.68
1.20
1.09
0.4
0.3
LiAsF0
LiPF0
a Contains 0.1M TBAFB.
El00 is that anodic or cathodic voltage where the current density
increases to 100 ~tA/cm2.
t
315
r,
2!lS
1-15
(J'.lS
POTENTIAL (VOLTS) VS L i / L i § (2,ZM)
Fig. 3. CV curves for LiAsF6 in 0 . I M TBAFB/THF at GC. Scan rate
100 mV/s, oreo ot GC = 0.07 cm 2. (A) 0.009M LiAsF6, and (B) 2.1M
LiAsF6.
=
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J. Electrochem. Soc.: E L E C T R O C H E M I C A L S C I E N C E A N D T E C H N O L O G Y
2916
2
5(
-4(
-3{
4
1
~-10
1
,/
4"6
470
3-0
2"0
1.0
0'6
POTE,T~AL VS U ~V0LTS)
Fig. 4. CV curves of 3.9 mM LiSbF6 in; 1M TBAFB/THF at GC,
- 1 st cycle, - ..... 2nd cycle. Scan rate = 100 mV/s, area of GC =
0.07 cm 2.
(peak 5) was seen at 3.7V (dotted curve). Addition of authentic SbF3 to the above solution and running a CV (Fig.
5a) reproduced the cathodic peak 4 and the corresponding
oxidation peak 5. Therefore, the reduction and oxidation
reaction at peaks 4 and 5, respectively, can be written as
SbF3 + 3e- ~Sb(0) + 3F-
[3]
Because the integration under the cathodic and anodic
peaks is identical (2.1 • 10 4C), there are no chemical follow-up reactions in the time scale of the scan.
We propose that the cathodic reaction under peaks 1 and
2 is the reduction of Sb(V) to Sb(0). However, Sb(0) is not
formed quantitatively since the anodic charge under peak
3 is only 2% of the cathodic charge under peaks 1 and 2.
The following reaction steps can be written to explain the
reduction occurring at peaks 1 and 2
Sb(V) + 5e- ~- Sb(0) + other Sb-containing products [4]
It is also instructive to consider how the position of the
Sb(0) stripping peak potential depends on the ionic envi-
,
l
I
-80
A
0
/
/
/
8O
160
3!o
4'5
31s
I
'
z'.7
'
211
1.5
POTENTIAL VS. Li (VOLTS)
J
J
I
-40
0
_
_
+F~
.........
9
f _ J - -
/
[ ,
o
401
80
J
4
p
3
i
2
POTENTIAL VS. L• (VOLTS)
Fig. 5. CV curve of (A) 5.4 mM SbFs, and (B) 11.2 mM SbFs and SbF3
in 1M TBAFB/THF. Scan rate = 100 rnV/s, area of GC = 0.07 cm2.
December 1988
r o n m en t at the electrode surface. When a metal dissolves,
the potential of dissolution is affected by the type of anion
and its concentration. It is known, for instance, that the
dissolution potential of Ag is shifted by -0.8V in the presence of C1- ion due to complexation (11).
At very high concentrations of LiSbF~, peaks 1 and 2
overlap and only a shoulder is seen instead of a distinct
peak 2. Peak potentials are also shifted an o d i c~ l y as the
concentration of LiSbF~ increases. As in the case of
LiAsF6, peak currents are m u c h lower than expected
based on the measured current at lower concentration (3.9
mM). It is possible that LiF formed as in Eq. [2] can passir a t e the electrode. The potential is shifted positive both
due to LiF formation and SbF~ reduction at peaks 1 and 2.
TO ascertain whether SbF5 was responsible for one of the
observed reduction peaks, a CV (Fig. 5b) was run in a solution containing 11.2 mM of SbF5 in THF. The SbF~ was intercalated on graphite and the species we detect electrochemically have diffused out of the solid into the bulk
electrolyte. While SbF~ intercalated graphite is k n o w n to
contain mostly (>90%) SbF~, low levels of SbF3 and S b F C
may also be present (12, 13). Essentially all of the SbFs is
weakly bound to the surface and readily washes off (14),
thus explaining the small SbF3 reduction peak in our solution. While some uncertainty remains, the prominent peak
at 1.0V is likely reduction of the major species present,
SbFs. Since the electrode is partially passivated at these
potentials, the exact peak potential for the reduction of
SbF5 cannot be compared with tha~t of peak 1 or peak 2 of
Fig. 4. It is likely that one of these peaks is the result of reduction of SbF5 formed by the following equilibrium
SbFC ~-- SbF~ + F-
[5]
The behavior of LiSbF6 in sulfolane is similar to that observed in THF except that in sulfolane there is an additional reduction peak at 0.6V. An anodic peak at 1.2V due
to the oxidation of the product generated at 0.6V was also
seen. This peak is observed only a~ very high concentrations of LiSbF8 (0.8M). The nature of this redox couple is
clearly solvent dependent and may well involve S b F J s u l folane reaction products.
Of all the anions investigated in this report, only SbFC
w a s observed to reduce on Hg. It gave two reduction
waves, one at 0.9V and the other at 0.65V. But, at concentrations greater than 10 3M, the waves were not well defined owing to polarographic maxima (15). Unlike SbF6 ,
A s F c is not observed to reduce on the Hg electrode. As reported earlier, AsF6- is reduced on a GC electrode. This
fact suggests that there is a kinetic inhibition for the reduction of A s F c at Hg:
Factors governing reactivity.--Why do these anions behave differently? It is necessary to assess both chemical
and electrochemical stability, both of which are required
of an ideal anion. We propose that chemical stability of
these anions is most conveniently understood by considering factors governing the Lewis acidity of the MF5
species. The electrochemical stability is likely to be both
kinetic and thermodynamic in nature. It can be understood in terms of the degree of shielding and the effective
nuclear charge seen by an electron at the surface of the Li
electrode as it interacts with the MF6- anion.
While the factors responsible for the relative ease of
S b F C reduction cannot be stated with certainty, the decreasing charge to volume ratio (i.e., P F C > A s F c >
SbF6-) which occurs as one descends the Group VA colu m n is noteworthy. The spreading of an identical amount
o f outer shell negative charge over larger surface regions
leads to poorer shielding of the central ion in the larger
MF6- species and greater electron affinity.
The above experimental and theoretical considerations
point to PF6 as the optimum anion from an electrochemical viewpoint. But PF6 has a major deficiency, and AsF6to a lesser extent, when one considers chemical reactivity,
particularly in the presence of Li +.
We propose that the major property which governs the
chemical stability of MFC anions in the presence of Li + is
the Lewis acid strength of the conjugate MF~ species. The
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VoL. 135, No. 12
LiMF6 I N T E T R A H Y D R O F U R A N A N D S U L F O L A N E
stronger pentafluoro Lewis acids, MFs, will more effectively bind the Lewis base, F , and form a more chemically
stable anion MFC. A great deal of experimental evidence
indicates the order of increasing Lewis acidity to be PF5 <
AsF~ < SbF~ (16). Therefore, SbF6 is likely the most chemically stable anion, at least with regard to loss of F- and
Lewis acid formation. The presence of MF~ in electrolytes,
whether PFs, AsFs, SbFs, or others, is well k n o w n to polymerize and degrade ether solvents. Thus, very low MF5 levels must be maintained for long-term electrolyte stability.
In the presence of most cations, even Na +, all three MF6anions appear to resist F- dissociation quite well. However, the extremely high crystal lattice energy of LiF, in
conjunction with LiF insolubility in aprotic solvents,
drives the loss of F- from MFC, resulting in the following
overall equilibrium
Li + + MF6- ~- MF5 + LiF
[6]
Unfortunately, the previously discussed factors which
favor good electrochemical stability lead to increased reactivity in the presence of Li + counterion.
Summary and Conclusions
This investigation has shed light on our understanding
of t h e reduction mechanisms of SbF~- and AsF~ . MF6anions were found to reduce in the order Sb > AS > P
which closely parallels a decreasing core ion radius and increasing charge to volume ratio. The nonreducibility of
P F ( indicates that it is an electrochemically stable salt for
cycling the Li" electrode, but its disproporti0nation to PF~
precludes its use as a stable anion in rechargeable Li batteries. The extreme Lewis acidity of SbF5 confers good
chemical stability on LiSbF6, but results in an electrochemically inferior salt. Accordingly, AsF6- is still the best
Li + counterion due to its u n i q u e ability to compromise between ease of reduction to M(III) and M(0) and disproportionation to MFs.
Acknowledgment
This work was supported by the National Science Foundation, and the National Aeronautics and Space Administration.
2917
Manuscript submitted Jan. 22, 1988; revised manuscript
received April 21, 1988. This was Paper 75 presented at the
Honolulu, HI, Meeting of the Society, October 18-23, 1987.
REFERENCES
1. K. M. Abraham, J. L. Goldman, and D. L. Natwig, This
Journal, 129, 2404 (1982).
2. V. R. Koch, ibid., 126, 181 (1979).
3. R. J. H o m i n g and W. B. Ebner, U.S. Pat. 4,107,404
(1978).
4. Y. Matsuda, Electrochemical Society of Japan Extended Abstracts, "Third International Meeting on
Lithium Batteries," p. 322, Kyoto, Japan, May 27-30,
1986.
5. D. T. Sawyer and J. L. Roberts, Jr., in "Experimental
Electrochemistry for Chemists," p. 226, John Wiley
& Sons, Inc., New York (1974).
6. N. Margalit and H. J. Canning, in "Proceedings of the
Electrochemical Society, Battery Division," Vol.
80-4, B.B. Owens and N. Margalit, Editors, p. 334,
The Electrochemical Society Softbound Proceedings Series, Princeton, NJ (1980).
7. A. N. Dey and E. J. Rudd, This Journal, 121, 1294
(1974).
8. J. L. Goldman, R. M. Mank, J. H. Young, and V. R.
Koch, ibid., 127, 1461 (1980).
9. R. A. Wiesbock, U.S. Pat. 3,654,330 (1972).
10. A. J. Bard and L. R. Faulkner, in "Electrochemical
Methods," p. 24, John Wiley & Sons, Inc., New York
(1980).
11. J. B. Lambling and G. Cauquis, in "Advances in Analytical Chemistry and Instrumentation," Vol. 10,
H. W. Nurenberg, Editor, p. 380, John Wiley & Sons,
Inc., New York (1974).
12. L. Facchini, J. Bovat, H. Sfihi, A. P. Legrand, G. Furdin, J. Melin, and R: Vangelisti, Synthetic Metals, 5,
11 (1982).
13. G. Roth, K. Lfiders, and H.J. Gfintherodt, ibid., 8
(1983).
14. W. C. Forsman, T. Dziemianowicz, K. Leong, and D.
Carl, ibid., 5, 77 (1983).
15. A. J. Bard and L. R. Faulkner, in "Electrochemical
Methods," p. 151, John Wiley & Sons, Inc., New
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(1985).
Characterization of Li/SO2CI2 and Li/"SO2CI2 + CI2" Cells
K. M. Abraham,* M. Alamgir,* and R. K. Reynolds
EIC Laboratories, Incorporated, Norwood, Massachusetts 02062
ABSTRACT
A comparative study of the discharge performance of LYSO~Clz and Li/"SO~C12 + C12" cells was made. The low temperature discharge capacities of both types of cells appear to be influenced by the complexes formed between Li § and the
solvents. At -20~ better capacities were obtained at the higher discharge rate of 1.0 than at 0.2A. Both of these cells exhibited a tendency for safety hazards during forced overdischarge. The overdischarge-related safety features of the cells
appeared to be dependent on the ratio of the active components and may be improved by appropriate design considerations.
The development of the Li/SO2C12 cell has been pursued
because of its ability to offer a n u m b e r of potential advantages over the Li/SOC12 cell (1-6). A popular belief has been
that the Li/SO2C12 cell is safe because of the absence of S as
a discharge product. Another advantage of the Li/SO2C12
cell is its higher open-circuit potential of 3.90V which
should translate into higher load voltages. Although the
expected high load voltage and capacity were realized at
low rates, these advantages did not persist at high discharge rates, due to excessive cathode polarization. Solutions to the cathode polarization problem were found with
the use of higher surface area carbon cathodes (3) and by
carbon catalysis with Pt (2) and transition metal phthalo* Electrochemical Society Active Member.
cyanines (6). The Li/SO2C12 cell has been found to suffer
from excessive self-discharge because of a higher rate of Li
anode corrosion than that experienced in the Li/SOC12 cell.
It has apparently been possible to mitigate anode corrosion by reverse polarity cell design, use of alternate electrolyte salts, and by addition of a cosolvent such as SO2 (2).
I n an attempt to impro;ee the performance of the
Li/SO2C12 cell, Liang and co-workers (7-10) used -0.5M C12
as an additive to the SO2C1j1M LiA1C14 electrolyte. These
Li/SO2C12 cells containing C12, referred to as CSC cells, exhibited high discharge rate capability. More than one CSC
cell design was used to maximize performance at high and
low temperature regimes encompassing -30 ~ to 150~ (8,
10).
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