12 Chemical Investigations of Lithium-Sulfur Cells J. R. BIRK and R. K. S T E U N E N B E R G Downloaded by UNIV OF ARIZONA on December 20, 2012 | http://pubs.acs.org Publication Date: September 1, 1975 | doi: 10.1021/ba-1975-0140.ch012
1
Chemical Engineering Division, Argonne National Laboratory, Argonne, Ill. 60439
High-performance lithium-sulfur cells are being developed for use in off-peak energy storage batteries in electric utility networks. The cells, which operate at 400°C, consist of a lithium electrode, a sulfur or sulfide electrode, and molten LiCl-KCl electrolyte. The chemistry of the lithium electrode is relatively straightforward. However, the electrochemical reactions at the sulfur electrode involve the formation of several intermediate species that are sufficiently soluble in the electrolyte to limit the lifetime and capacity of the cells. Although this effect can be decreased with soluble additives such as arsenic or selenium in the sulfur, a more promising solution appears to be the use of metal sulfides, rather than sulfur, as the active material in the positive electrode.
TTigh-performance lithium-sulfur secondary batteries are being developed for use in electric automobiles and for off-peak energy storage in electric utility systems. These applications impose severe performance requirements that cannot be met by present batteries. For the electric automobile, the battery must have a minimum specific energy of >—'200 W-hr/kg, a specific power of at least 200 W/kg, and a lifetime of 3-5 yrs. The projected performance requirements for a battery for off-peak energy storage are a maximum cost of $12-15/kW-hr, a specific power of /—50 W/kg, and a minimum lifetime of 5-10 yrs. The development of lithium-sulfur batteries began in 1967 as a small basic research effort at Argonne National Laboratory (I). Since then, 'Present address: Electric Power Research Institute, P. O. Box 10412, Palo Alto, Calif. 94304. 186 In New Uses of Sulfur; West, J.; Advances in Chemistry; American Chemical Society: Washington, DC, 1975.
Downloaded by UNIV OF ARIZONA on December 20, 2012 | http://pubs.acs.org Publication Date: September 1, 1975 | doi: 10.1021/ba-1975-0140.ch012
12.
Lithium-Sulfur
BIRK A N D S T E U N E N B E R G
Cells
187
significant progress has been made both at Argonne and at other laboratories, including Atomics International and General Motors. A t present, the Argonne program is entering the early hardware stage, with the construction and testing of a full-scale sealed cell of 120 A - h r (156 W - h r ) capacity and a specific energy of 104 W-hr/kg (2). Recent developments in both hardware and cell life indicate that a lithium-sulfur battery system capable of meeting the above requirements may be forthcoming within 5 yrs. Subsequent commercial development and production of the batteries may require another few years, thereby placing large-scale practical application in the early 1980s. The general aspects of the development work on lithium-sulfur batteries have been described in various technical papers and progress reports. This particular paper is concerned mainly with the chemical and electrochemical processes occurring in the cells. In addition, recommendations that have resulted in improved cell performance, some of which were derived as a result of this investigation, w i l l be discussed and evaluated. Lithium and sulfur are promising active electrode materials for batteries because of their low equivalent weight, low cost, and suitable electrochemical properties. The major question in the use of these active electrode materials is whether the electrodes will be able to provide sustained high performance. U n t i l recently, capacity retention over an extended period has been difficult to achieve. Chemistry of the Lithium Electrode The capacity loss from the lithium electrode has resulted mainly from dewetting of metallic substrate materials, such as nickel and stainless steel Feltmetal and reaction with the electrolyte, the L i C l - K C l eutectic, to form potassium vapor: Li(l)
+ KC1(1) -> K(g) +
LiCl(l)
(1)
This reaction results in an equilibrium potassium vapor pressure (calculated from thermodynamic data) of 0.714 ton* above the L i C l - K C l eutectic at 427°C ( 7 0 0 ° K ) . Metallic lithium is rapidly lost, by Reaction 1, from the lithium electrode in open cells exposed to an inert atmosphere of helium (3). However, this reaction has not been evident in hermetically sealed cells. Lithium is soluble in the L i C l - K C l electrolyte at about 0.13 mole % at 400°C, based on an extrapolation of the data from the literature (4, 5). This solubility does not lead to significant lithium loss. However, the electronic conductivity of the lithium-saturated salt at 450°C is estimated (6, 7) to be 5 X 10" ohm" c m , which could lead to extensive self-dis3
1
1
In New Uses of Sulfur; West, J.; Advances in Chemistry; American Chemical Society: Washington, DC, 1975.
Downloaded by UNIV OF ARIZONA on December 20, 2012 | http://pubs.acs.org Publication Date: September 1, 1975 | doi: 10.1021/ba-1975-0140.ch012
188
NEW
USES
O F SULFUR
charge. Fortunately, self-discharge rates of lithium-chlorine (8) and lithium-sulfur cells are much lower than that predicted from the electronic conductivity of the electrolyte saturated with lithium ( 8 ) . It is possible that oxidation of dissolved lithium at the sulfur electrode depletes the lithium concentration at the electrode surface, thereby causing a break in the electronic circuit within the electrolyte. The dewetting problem of the lithium electrode is now under intensive investigation. General Motors has found that high-temperature hydrogen treatment of the nickel Feltmetal substrate improves lithium retention ( 9 ) . In this process, metal-oxide films are removed from the substrate metal and the lithium and substrate are in more intimate contact. In this laboratory, solid lithium—aluminum alloys and solutions of lithium and metal additives are being tested. This latter approach takes advantage of the fact that solubility and wettability are related (JO) and that certain metal additives are somewhat soluble i n lithium at 400°C. As a result, good electrical and physical contact can be maintained between the lithium and metal additive and between the metal additive and substrate metal. Chemistry of the Sulfur Electrode The capacity loss of the sulfur electrode can be attributed to sulfur vaporization (the vapor pressure of sulfur at 400°C is 410 torr), migration or dispersion of insoluble sulfur-containing phases from the electrode, solubilization of sulfur-containing species i n the electrolyte, and/or i n activation of sulfur within the electrode compartment. Since neither the mechanism of the cell reaction nor the mechanism of sulfur loss was understood, a study of sulfur electrode chemistry and electrochemistry was made. It was expected that information gained from these studies would lead to improved performance and lifetimes of lithium-sulfur cells. Cyclic Voltametric Studies. The major investigative tool used to study the chemistry and electrochemistry of the sulfur electrode was cyclic voltammetry. The working solution contained L i S ( S -
(2)
Downloaded by UNIV OF ARIZONA on December 20, 2012 | http://pubs.acs.org Publication Date: September 1, 1975 | doi: 10.1021/ba-1975-0140.ch012
2
In arriving at this reaction, the authors assumed the identity of the product ( L i S ) , which they generated electrochemically, in order to prove their overall reaction. In addition, the concentrations of L i S , which they indicated were generated in the L i C l - K C l eutectic, were much above the solubility limit (0.029M at 4 2 5 ° C ) as determined by L i u et al. (16). Bernard et al. (12, 13) used chronopotentiometry and spectrophotometry to demonstrate that two intermediates, a polysulfide (S ~) and a supersulfide (S "), are involved i n the sulfide oxidation. A definitive reaction mechanism, however, could not be proposed because the authors stated (12) that their reaction was spontaneous and indicated (13) that they could not identify the oxidizing agent. Recent cyclic voltammetric studies by Kennedy and Adamo (14) and by Cleaver et al. (15) indicate that sulfur reduction involves two steps. These authors observed two peaks on both anodic and cathodic scans which they concluded were due to the two following steps: 2
2
w
2
n
nS
t± S
n
"
l± S
n
2
-
(3)
Neither group of workers indicated or mentioned a further reduction (e.g., Reaction 4 ) : S - + n
2
(2n—2)e~
