Lithium Batteries: A Practical Application of Chemical Principles

Chemistry for Everyone

Lithium Batteries: A Practical Application of Chemical Principles Richard S. Treptow Department of Chemistry and Physics, Chicago State University, Chicago, IL 60628-1598

Cameras, laptop computers, cellular telephones, and automobiles have one thing in common—they all require batteries. The devices of modern technology make ever-increasing demands for sources of portable energy. The ideal battery for every purpose would be dependable, economical, lightweight, stable in storage, and environmentally safe. Although the perfect battery does not exist, and perhaps never will, electrochemical engineers continue to invent new varieties with improved performance (1–7). In recent years new batteries have emerged in the marketplace that make use of the unique properties of lithium, the lightest metal on the periodic table. A collection of these commercial batteries is displayed in Figure 1. They differ in their size, shape, and voltage, and also in the chemical reactions they employ. Lithium batteries are of interest to chemistry students because of the increasing role they play in modern technology and also because they illustrate the practical application of principles learned in the classroom. In this paper we imagine ourselves in the role of engineers as we design lithium batteries through the application of these principles. We will discuss batteries whose anodes are lithium metal and also those that use lithium-ion anodes.

Figure 1. Commercial lithium-metal batteries of various sizes, shapes, voltages, and chemistries.

The Lithium-Metal Anode Metals have properties that make them well suited to serve as battery anodes. They are easily oxidized from their metallic state to produce ions and electrons, as illustrated by the half-reaction

M

Mn + + ne−

The electrons liberated are conducted through the metal, which thereby becomes the negative terminal of the cell. The fact that metals are physically strong and easily fashioned into any desired shape adds to their attractiveness as anodes.

Metals commonly employed as anodes in commercial batteries are listed in Table 1. The tabulated properties give clues as to the ability of each metal to play this role. The redox couple identifies the ion formed in the oxidation halfreaction. The standard reduction potential is a measure of the ease with which that ion is reduced back to the metallic state. The electrochemical capacity of a metal is a practical property defined as the quantity of electric charge produced per gram of metal in the anode half-reaction. For example, in the case

Table 1. Metals Commonly Used as Battery Anodes Metal

Atomic Weight

Density/ (g cm3)

Redox Couple

Standard Reduction Potential/V

Electrochemical Capacity/(A h g1)

Lithium

6.94

0.53

Li+/Li

3.04

3.86

+

Sodium

22.99

0.97

Na /Na

2.71

1.17

Magnesium

24.31

1.74

Mg2+/Mg

2.70

2.21

Aluminum

26.98

2.70

Al3+/Al

1.66

2.98

Calcium

40.08

1.54

Ca2+/Ca

2.87

1.34

Iron

55.85

7.86

Fe2+/Fe

0.45

0.96

Zinc

65.39

7.14

Zn2+/Zn

0.76

0.82

Cadmium

112.41

8.64

Cd2+/Cd

0.40

0.48

Lead

207.2

Pb2+/Pb

0.13

0.26

11.3

JChemEd.chem.wisc.edu • Vol. 80 No. 9 September 2003 • Journal of Chemical Education

1015


of lithium the electrochemical capacity is calculated as follows: 1 mol e− mol Li

×

mol Li 6.94 g Li

×

e−

9.65 × 104 C

–

mol e −

+

anode

×

1As C

×

h 3600 s

= 3.86

Ah g Li

The unit used to measure charge is the ampere-hour. Two properties of lithium listed in Table 1 suggest this metal is an attractive choice to serve as a battery anode. First, the lithium ion has the most negative reduction potential. This means that the ion has the lowest tendency to be reduced to the metallic state. More importantly, it means that lithium metal has the greatest tendency to be oxidized to the ionic state. If all other factors were equal, a lithium anode would produce a higher voltage than any of the other metals. Second, lithium has the greatest electrochemical capacity. As the calculation above shows, this results because lithium has an exceptionally low atomic weight. Thus, a lithium anode will deliver the greatest quantity of charge per unit-weight. Since the energy produced by a battery is the product of its voltage times the charge it delivers, the two properties cited suggest that lithium batteries can be expected to produce exceptionally high energy per unit-weight. A battery that we shall design employing a lithium anode is illustrated in Figure 2. When the cell discharges, the oxidation half-reaction at the anode is Li(s)

Li+(solv) + e−

The electrons produced leave the battery, pass through the external circuit with a load of some sort, and return to the battery at the cathode. The solvated lithium ions migrate through the cell to the cathode where they undergo a chemical reaction yet to be chosen. Let us first select the electrolyte through which this ion migration occurs. Electrolyte Salts and Solvents The role of a battery electrolyte is to separate the anode and cathode so they do not react directly. It must, however, allow ions to pass between the electrodes. This movement of ions constitutes the flow of charge inside the battery. For our lithium battery the electrolyte must allow lithium ions to migrate from the anode to the cathode. Most batteries use electrolytes that are aqueous solutions of salts or other ionic compounds. A lithium battery, however, must avoid water and other protic solvents because these compounds react with the metal to produce hydrogen gas. The battery must employ a solvent that is aprotic and nonreactive toward the metal. The solvent must be sufficiently polar to dissolve lithium salts and produce a solution with high ionic conductivity. Many organic solvents meet these criteria. They include acetonitrile (AN), dioxane (DIOX), γbutyrolactone (BL), methyl formate (MF), 1,2dimethoxyethane (DME), propylene carbonate (PC), dimethylsulfoxide (DMSO), and tetrahydrofuran (THF).

1016

load

cathode

electrolyte

FeS2 +4Li + + 4e−

Li Li + Li + + e−

Fe + 2Li2S

lithium metal

iron disulfide

Figure 2. Lithium-metal battery with a solid cathode of iron disulfide.

Our next task is to find lithium salts that are soluble in these solvents. In general, salts achieve high solubility if their lattice energies are low. The small lithium ion accomplishes this if it is paired with a large anion. Examples of such salts include LiClO4, LiPF6, LiBr, LiBF4, LiCF3SO3, and LiAlCl4. These salts are soluble in the solvents above, and the solutions they produce have the required high ionic conductivity. Suppose we choose a solution of LiPF6 in dimethylsulfoxide as the electrolyte for the battery shown in Figure 2. The final step is to select its cathode. Solid, Liquid, and Gas Cathodes The cathode of a lithium battery must undergo a reduction reaction that consumes the electrons and lithium ions produced at the anode. Iron disulfide is one of many inorganic solids that can serve in this role. Its reduction half-reaction is FeS2(s) + 4Li+(solv) + 4e−

Fe(s) + 2Li2S(s)

Let us choose this compound as the cathode of the battery in Figure 2. The overall cell reaction is then 4Li(s) + FeS2(s)

Fe(s) + 2 Li2S(s)

and the shorthand notation for the cell is Li(s)LiPF6, DMSOFeS2(s) Chemical thermodynamics can be applied to calculate the voltage of the battery. The process is simplified by the fact that all reactants and products in the overall reaction are solids. Thus, they are present in their standard states. The standard free-energy change for the reaction can be calculated from the standard free energies of formation by use of the equation ∆G

o

=

∑ n ∆Gfo (products)

−

∑ m ∆Gfo (reactants)

Journal of Chemical Education • Vol. 80 No. 9 September 2003 • JChemEd.chem.wisc.edu


e−

cathode its reduction half-reaction is

load

+ 2 SOCl2(l) + 4 Li (solv) + 4 e−

+

S(solv) + SO2(solv) + 4LiCl(s)

–

layer

Li

2SOCl2 + 4 Li + + 4e−

passive

Li + Li+ + e−

current collector

cathode

anode

S + SO2 + 4LiCl

lithium metal

thionyl chloride

Figure 3. Lithium-metal battery with liquid cathode of thionyl chloride.

Substituting the appropriate ∆G f values from the literature gives ∆G = 718.1 kJ. We then calculate the cell potential from o

Ecell

∆G nF

= −

o

The sulfur and sulfur dioxide produced remain dissolved in the thionyl chloride, but the lithium chloride precipitates. A battery we might build using this chemistry is illustrated in Figure 3. The lithium chloride precipitate plays an important role. It deposits on the surface of the lithium anode and thereby forms a passive layer between the anode and cathode. This layer prevents the anode and cathode from reacting directly and yet it allows the lithium ions produced at the anode to pass through. The lithium chloride plays the role given to the electrolyte phase in most other batteries. Next, we dissolve LiAlCl4 into the thionyl chloride to increase its ionic conductivity. Finally, a current collector of carbon or some other electronically conducting material is placed in contact with the cathode. The cell is now complete; its overall reaction is

4Li(s) + 2SOCl2(l)

S(s) + SO2(solv) + 4LiCl(s)

and its notation is

where F is the Faraday constant and n = 4 mol for this reaction. The result is E cell = 1.86 V. Electrochemical engineers refer to this as the open-circuit voltage. It is the potential between the terminals of a battery when it is not delivering a current. When a battery is delivering a current its voltage decreases as a result of cell-polarization effects (1, 4, 8). Commercial batteries that employ the FeS2 cathode are rated at an operating voltage of 1.5 V. A few inorganic liquids are able to serve as both the cathode and the electrolyte solvent of a lithium battery. Thionyl chloride is the most common example. In the role of the

Li(s)LiAlCl4, SOCl2(l)C(s)

Applying the calculation method described above we find that E cell = 3.71 V for this battery. The calculation must be regarded as approximate, however, when applied to any cell with a liquid cathode because the reactants and products in the liquid phase are not in their standard states. Commercial batteries based on this chemistry are called inorganic lithium batteries because no organic solvent is present. They have an operating voltage of 3.6 V. Various solid and liquid compounds employed as cathodes in lithium batteries are listed in Table 2. The table gives

Table 2. Solid and Liquid Cathodes for Lithium Batteries Cathode Compound

Cell Reaction

E cell / V

Solids Oxides CuO

2Li(s) + CuO(s) → Cu(s) + Li2O(s)

2.25

Bi2O3

6Li(s) + Bi2O3(s) → 2Bi(s) + 3Li2O(s)

2.06

Sulfides FeS

2Li(s) + FeS(s) → Fe(s) + Li2S(s)

1.75

FeS2

4Li(s) + FeS2(s) → Fe(s) + 2Li2S(s)

1.86

Ni3S2

4Li(s) + Ni3S2(s) → 3Ni(s) + 2Li2S(s)

1.73

CuS

2Li(s) + CuS(s) → Cu(s) + Li2S(s)

2.00

Halides CuF2

2Li(s) + CuF2(s) → Cu(s) + 2LiF(s)

3.55

CuCl2

2Li(s) + CuCl2(s) → Cu(s) + 2LiCl(s)

3.08

AgCl

Li(s) + AgCl(s) → Ag(s) + LiCl(s)

2.84

SOCl2

4Li(s) + 2SOCl2(l) → S(solv) + SO2(solv) + 4LiCl(s)

3.71

SO2Cl2

2Li(s) + SO2Cl2(l) → SO2(solv) + 2LiCl(s)

3.90

Liquids


1017


the overall reaction for the cell in which each compound serves as the cathode, and it lists E cell calculated from theory as discussed above. The ∆G f values at 25 C needed to calculate these cell potentials are given in Table 3 (9). Another class of lithium-metal batteries uses dissolved sulfur dioxide gas as the cathodic reactant. The cell reaction is 2 Li(s) + 2 SO2(solv)

Li2S2O4(s)

The sulfur dioxide is dissolved in an electrolyte solution such as a solution of LiBr in acetonitrile. The lithium dithionite precipitate forms a passive layer on the anode preventing it from reacting directly with the sulfur dioxide. As with the battery shown in Figure 3, a current collector is placed in contact with the liquid cathode. The resulting battery would be symbolized

Table 3. ∆Gf° for Selected Compounds in Lithium Batteries Compound

∆Gf°/ (kJ mol1) 109.77

AgCl(s)

∆Gf°/ (kJ mol1)

Compound LiF(s)

588.66

Bi2O3(s)

493.47

LiCl(s)

384.02

CuF2(s)

491.64

Li2O(s)

562.11

CuCl2(s)

173.81

Li2S(s)

439.08

CuO(s)

128.29

Ni3S2(s)

210.39

SO2(g)

300.10

CuS(s)

53.47

FeS(s)

101.96

SOCl2(l)

202.61

FeS2(s)

160.07

SO2Cl2(l)

315.30

Li(s)LiBr, SO2, ANC(s) e−

Commercial batteries employing the SO2 cathode have an operating voltage of 2.8 V.

+ – anode

Some inorganic solids have the ability to serve as cathodes for lithium batteries through a unique type of reaction known as intercalation. These solids have layered or tunneled crystal structures. In the cathode half-reaction, the crystal acts as a host for lithium ions that diffuse into it, as illustrated in the battery shown in Figure 4. The host is reduced by the gain of electrons required for electro-neutrality as the insertion occurs. The host undergoes only minor structural change in the process. The half-reaction can be written as LinHOST(s)

The intercalation compound formed in the process is nonstoichiometric. The most popular intercalation cathode for commercial lithium-metal batteries is manganese dioxide. Its reduction half-reaction is MnO2(s) + nLi+(solv) + ne−

LinMnO2(s)

and its overall cell reaction is

nLi(s) + MnO2(s)

LinMnO2(s)

The cathode consists of manganese dioxide with carbon particles added for improved conductivity. A suitable electrolyte is a solution of LiClO4 in propylene carbonate (PC). With these components the cell shown in Figure 4 is symbolized as

Li(s)LiClO4, PCMnO2(s)C(s) Commercial versions of this battery are rated at 3.0 V or at multiples of this voltage if they are composed of two or more individual cells connected in series. Several intercalation hosts used as cathodes for lithium batteries are listed in Table 4. The cell reaction is shown for each cathode. The Ecell listed is the measured open-circuit 1018

cathode

electrolyte

Li Li +

Li+

Li+

Li + + e−

current collector

Intercalation Cathodes

HOST(s) + n Li+(solv) + n e−

load

manganese dioxide and carbon

lithium metal

Figure 4. Lithium-metal battery with intercalation cathode of manganese dioxide.

Table 4. Intercalation Hosts Used as Cathodes for Lithium Batteries Host Compound

Cell Reaction

Ecell/V

CoO2

nLi(s) + CoO2(s) → LinCoO2(s)

3.8–4.5

V2O5

nLi(s) + V2O5 (s) → LinV2O5(s)

2.9–3.6

MnO2

nLi(s) + MnO2(s) → LinMnO2(s)

2.5–3.2

TiS2

nLi(s) + TiS2 (s)

MoO2

nLi(s) + MoO2(s) → LinMoO2(s)

→ LinTiS2(s)

1.6–2.5 0.8–1.9

voltage of the cell (1). It has a range of values because the voltage changes with n, the extent of intercalation of the host. The Lithium-Ion Anode The reactions discussed above in which lithium ions diffuse into a host structure have the advantage of being reversible. Hence, they can be used to produce rechargeable



current collector

–

e−

load or dc power source

discharge

Li

charge

Li +

+

Li

+

discharge

Li

+

Li

+

Li+

charge lithiated graphite

electrolyte

+

current collector

e−

cobalt dioxide

Figure 5. Lithium-ion rechargeable battery with intercalation positive and negative electrodes during discharge and charge.

called a lithium-ion battery to emphasize that it contains no lithium metal. The electrode that serves as the cathode during discharge of the battery in Figure 5 is cobalt dioxide. Its reduction halfreaction is CoO2(s) + n Li+(solv) + n e−

LinCoO2(s)

Thus, the overall discharge reaction is LinC6(s) + CoO2(s)

6 C(s) + LinCoO2(s)

This reaction occurs in the opposite direction when the battery is charged. The battery requires an electrolyte, conductive carbon particles in the CoO2(s), and current collectors as described above. Commercial versions have an operating voltage of 3.7 V. The rechargeable battery in both the discharging and charging modes is illustrated in Figure 5. When discharging the battery delivers energy to the external load and when charging it receives energy from a dc power source. The electrons and lithium ions reverse directions whenever the mode changes. Batteries of this type are called “swing” or “rocking chair” batteries to emphasize this back and forth motion. As with any rechargeable battery, use of the terms anode and cathode to describe electrodes can cause confusion. The electrode that acts as an anode during discharging becomes a cathode during charging. Electrochemical engineers avoid confusion by referring to the electrodes as positive and negative. These properties remain unchanged when the battery changes mode. Commercial Battery Construction

Figure 6. Interior of cylindrical lithium-ion rechargeable battery showing the spiral winding of electrodes and separators. Courtesy SANYO Energy (USA) Corporation.

batteries. The most commercially successful of these batteries is illustrated in Figure 5. Both electrodes are intercalation compounds. The electrode on the left serves as the anode when the cell discharges. It is a special intercalation compound consisting of a graphite host into which lithium ions have been electrochemically inserted between the carbon atom layers. The compound, known as lithiated graphite, is written LinC6. During discharge the lithium ions are extracted by the half-reaction

LinC6(s)

+

nLi (solv) + 6C(s) + ne−

This oxidation process is nearly as favorable as that discussed above for lithium metal. In fact, any battery that uses a lithiated-graphite anode will produce a voltage only about 0.2 V less than that of its lithium-metal anode counterpart. It is

The battery drawings featured above are schematic illustrations only. They do not show the many internal features needed to produce a practical battery. These features may include a gas-release vent to prevent the buildup of internal pressure. Separator layers of glass or polymeric fibers are commonly placed between the electrodes to prevent them from making direct contact. These separators must be permeable to the electrolyte solution. Finally, to deliver optimum current and power a practical battery must maximize the surface area of its electrodes. This is accomplished by fabricating the electrodes into thin sheets. The internal components of a commercial lithium-ion battery are shown in Figure 6. It is a cylindrical cell whose electrodes are thin sheets wound into a “jelly-roll” spiral with separator layers between them. Further information on commercial lithium batteries can be obtained by visiting the Web sites of companies that manufacture (10), market (11), or recycle (12) them. In closing, we should note that much current research is being done worldwide on the subject of lithium batteries. Their technology will surely change in years to come. Literature Cited 1. Handbook of Batteries, 3rd ed.; Linden, D., Reddy, T. B., Eds.; McGraw-Hill: New York, 2002. 2. Handbook of Battery Materials; Besenhard, J. O., Ed.; WileyVCH: Weinheim, Germany, 1999. 3. Lithium Ion Batteries; Wakihara, M., Yamamoto, O., Eds.; Wiley-VCH: Weinheim, Germany, 1998.


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Chemistry for Everyone 4. Vincent, C. A.; Scrosati, B. Modern Batteries: An Introduction to Electrochemical Power Sources, 2nd ed.; Arnold: London, 1997. 5. Julien, C. In The CRC Handbook of Solid State Electrochemistry; Gellings, P. J., Bouwmeester, H. J. M., Eds.; CRC Press: Boca Raton, FL, 1997. 6. Crompton, T. R. Battery Reference Book; Butterworths: London, 1990. 7. Lithium Battery Technology; Venkatasetty, H. V., Ed.; John Wiley & Sons: New York, 1984. 8. Treptow, R. S. J. Chem. Educ. 2002, 79, 334–338.

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9. HSC Chemistry for Windows, version 4.1; Outokumpu Research Oy: Pori, Finland, 2000. 10. Energizer. http://energizer.com. HDS Systems. http:// www.hdssystems.com. SANYO. http://sanyobatteries.com. Sonnenschein GmbH. http://www.sonnenschein-lithium.de. Valence Technology. http://www.valence-tech.com. Varta. http:// varta.com. Wilson Greatbatch Technologies. http:// www.greatbatch.com (all Web sites accessed Jun 2003). 11. Batteries Plus. http://batteriesplus.com. Interstate Batteries. http://interstatebatteries.com (all Web sites accessed Jun 2003). 12. ToxCo, Inc. http://toxco.com (accessed Jun 2003).


Lithium Batteries: A Practical Application of Chemical Principles

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