Batteries, Overview

Batteries, Overview

ELTON J. CAIRNSLawrence Berkeley National Laboratory and

University of California, Berkeley

Berkeley, California, United States

1. How Batteries Work

2. History

3. Types of Batteries and Their Characteristics

4. Applications

5. Environmental Issues

6. Future Outlook

Glossary

battery An assemblage of cells connected electrically inseries and/or parallel to provide the desired voltage andcurrent for a given application.

cell An electrochemical cell composed of a negativeelectrode, a positive electrode, an electrolyte, and acontainer or housing.

electrode An electronically conductive structure that pro-vides for an electrochemical reaction through thechange of oxidation state of a substance; may containor support the reactant or act as the site for theelectrochemical reaction.

electrolyte A material that provides electrical conductionby the motion of ions only; may be a liquid, solid,solution, polymer, mixture, or pure substance.

primary cell A cell that cannot be recharged and isdiscarded after it is discharged.

secondary cell A rechargeable cell.

Batteries are an important means of generating andstoring electrical energy. They are sold at a rate ofseveral billions of dollars per year worldwide. Theycan be found in nearly all motor vehicles (e.g.,automobiles, ships, aircraft), all types of portableelectronic equipment (e.g., cellular phones, compu-ters, portable radios), buildings (as backup powersupplies), cordless tools, flashlights, smoke alarms,heart pacemakers, biomedical instruments, wrist-watches, hearing aids, and the like. Batteries are souseful and ubiquitous, it is difficult to imagine howlife would be without them. Batteries, strictly speak-

ing, are composed of more than one electrochemicalcell. The electrochemical cell is the basic unit fromwhich batteries are built. A cell contains a negativeelectrode, a positive electrode, an electrolyte heldbetween the electrodes, and a container or housing.Cells may be electrically connected to one another toform the assembly called a battery. In contrast to thealternating current available in our homes from theelectric utility company, batteries deliver a directcurrent that always flows in one direction. There area few different types of batteries: Primary batteriescan be discharged only once and then are discarded;they cannot be recharged. Secondary batteries arerechargeable. Forcing current through the cells in thereverse direction can reverse the electrochemicalreactions that occur during discharge. Both primaryand secondary batteries can be categorized based onthe type of electrolyte they use: aqueous, organicsolvent, polymer, ceramic, molten salt, and so on.

1. HOW BATTERIES WORK

The electrical energy produced by batteries is theresult of spontaneous electrochemical reactions. Thedriving force for the reactions is the Gibbs freeenergy of the reaction, which can be calculated easilyfrom data in tables of thermodynamic properties.The maximum voltage that a cell can produce iscalculated from this simple relationship:

E ¼ �DG=nF; ð1Þwhere E is the cell voltage, DG is the Gibbs freeenergy change for the cell reaction (joules/mole), n isthe number of electrons involved in the reaction(equivalents/mole), and F is the Faraday constant(96487 coulombs/equivalent).

It is clear from Eq. (1) that a large negative valuefor the Gibbs free energy of the cell reaction isdesirable if we wish to have a significant cell voltage.

B

Encyclopedia of Energy, Volume 1. r 2004 Elsevier Inc. All rights reserved. 117

In practical terms, cell voltages of 1 to 2 V areachieved when using aqueous electrolytes, and cellvoltages up to approximately 4 V are achieved whenusing nonaqueous electrolytes. When larger voltagesare required, it is necessary to connect a number ofcells electrically in series. For example, 12-volt auto-motive batteries are composed of six 2-volt cellsconnected in series.

An important property of batteries is the amountof energy they can store per unit mass. This is thespecific energy of the battery, usually expressed inunits of watt-hours of energy per kilogram of batterymass (Wh/kg). The maximum value of the specificenergy is that which can be obtained from a certainmass of reactant materials, assuming the case thatany excess electrolyte and terminals have negligiblemass. This is called the theoretical specific energyand is given by this expression:

Theoretical Specific Energy ¼ �DG=SMw; ð2Þwhere the denominator is the summation of themolecular weights of the reactants.

It can be seen from Eq. (2) that the theoreticalspecific energy is maximized by having a largenegative value for DG and a small value for SMw.The value of DG can be made large by selecting forthe negative electrode those reactant materials thatgive up electrons very readily. The elements withsuch properties are located on the left-hand side ofthe periodic chart of the elements. Correspondingly,the positive electrode reactant materials shouldreadily accept electrons. Elements of this type arelocated on the right-hand side of the periodic chart.Those elements with a low equivalent weight arelocated toward the top of the periodic chart. Theseare useful guidelines for selecting electrode materials,and they help us to understand the wide interest inlithium-based batteries now used in portable electro-nics. Equations (1) and (2) give us a useful frame-work for representing the theoretical specific energyand the cell voltage for a wide range of batteries, asshown in Fig. 1. The individual points on the plotwere calculated from the thermodynamic data. Thesepoints represent theoretical maximum values. Thepractical values of specific energy that are availablein commercial cells are in the range of one-fifth toone-third of the theoretical values. As expected, thecells using high equivalent weight materials (e.g.,lead, lead dioxide) have low specific energies,whereas those using low equivalent weight materials(e.g. lithium, sulfur) have high specific energies. Thelines in Fig. 1 represent the cell voltages and simplyrepresent the relationships given by Eqs. (1) and (2).

The details of the electrochemical reactions in cellsvary, but the principles are common to all. Duringthe discharge process, an electrochemical oxidationreaction takes place at the negative electrode. Thenegative electrode reactant (e.g., zinc) gives upelectrons that flow into the electrical circuit wherethey do work. The negative electrode reactant is thenin its oxidized form (e.g., ZnO). Simultaneously, thepositive electrode reactant undergoes a reductionreaction, taking on the electrons that have passedthrough the electrical circuit from the negativeelectrode (e.g., MnO2 is converted to MnOOH). Ifthe electrical circuit is opened, the electrons cannotflow and the reactions stop.

The electrode reactions discussed in the precedingcan be written as follows:

Zn þ 2OH� ¼ ZnO þ H2O þ 2e�

and

2MnO2 þ 2H2O þ 2e� ¼ 2MnOOH þ 2OH�:

The sum of these electrode reactions is the overallcell reaction:

Zn þ 2MnO2 þ H2O ¼ ZnO þ 2MnOOH:

Notice that there is no net production or consump-tion of electrons. The electrode reactions balanceexactly in terms of the electrons released by thenegative electrode being taken on by the positiveelectrode.

100

1,000

10,000

0.01 0.1 1

The

oret

ical

spe

cific

ene

rgy

(Wh/

kg)

4.0 Volts

3.0

Pb/PbO2

MH/NiOOH

Zn/NiOOH

LiC /V ONa/S

Li/S

LiC 6 /Mn 2O4

LiC6 /CoO2

Zn/Air

Li/FeS 2

6 6 132.0

1.5

1.0

Equivalent weight (kg/equiv.)

FIGURE 1 Theoretical specific energy for various cells as a

function of the equivalent weights of the reactants and the cell

voltage.

118 Batteries, Overview

2. HISTORY

Volta’s invention of the Volta pile in 1800 representsthe beginning of the field of battery science andengineering. The pile consisted of alternating layersof zinc, electrolyte soaked into cardboard or leather,and silver. Following Volta’s report, many investiga-tors constructed electrochemical cells for producingand storing electrical energy. For the first time,relatively large currents at high voltages wereavailable for significant periods of time. Variousversions of the pile were widely used. Unfortunately,there was a lack of understanding of how the cellsfunctioned, but as we now know, the zinc waselectrochemically oxidized and the native oxide layeron the silver was reduced. The cells were ‘‘recharged’’by disassembling them and exposing the silverelectrodes to air, which reoxidized them. Inevitably,many other electrochemical cells and batteries weredeveloped.

John F. Daniell developed a two-fluid cell in 1836.The negative electrode was amalgamated zinc, and thepositive electrode was copper. The arrangement of thecell is shown in Fig. 2. The copper electrodes wereplaced in (porous) porcelain jars, which were sur-rounded by cylindrical zinc electrodes and placed in alarger container. A copper sulfate solution was placedin the copper electrode’s compartment, and sulfuricacid was put in the zinc electrode compartment.

The electrode reactions were as follows:

Zn ¼ Znþ2 þ 2e�

and

Cuþ2 þ 2e� ¼ Cu:

Sir William R. Grove, a lawyer and inventor of thefuel cell, developed a two-electrolyte cell related tothe Daniell cell in 1839. Grove used fuming nitricacid at a platinum electrode (the positive) and zinc insulfuric acid (the negative). Variants on this formula-tion were popular for a number of years. Of course,all of these cells were primary cells in that they couldbe discharged only once and then had to bereconstructed with fresh materials.

Gaston Plante invented the first rechargeablebattery in 1860. It was composed of lead sheetelectrodes with a porous separator between them,spirally wound into a cylindrical configuration. Theelectrolyte was sulfuric acid. These cells displayed avoltage of 2.0 V, an attractive value. During the firstcharging cycles, the positive electrode became coatedwith a layer of PbO2. The charging operation wascarried out using primary batteries—a laboriousprocess. Because of the low cost and the ruggednessof these batteries, they remain in widespread usetoday, with various evolutionary design refinements.

The electrode reactions during discharge of thelead–acid cell are as follows:

Pb þ H2SO4 ¼ PbSO4 þ 2Hþ þ 2e�

and

PbO2 þ H2SO4 þ 2Hþ þ 2e� ¼ PbSO4 þ 2H2O:

Notice that sulfuric acid is consumed during dis-charge and that the electrolyte becomes more dilute(and less dense). This forms the basis for determiningthe state of charge of the battery by measuring thespecific gravity of the electrolyte. The theoreticalspecific energy for this cell is 175 Wh/kg, a rather lowvalue compared with those of other cells.

Waldemar Jungner spent much of his adult lifeexperimenting with various electrode materials inalkaline electrolytes. He was particularly interested inalkaline electrolytes because there generally was nonet consumption of the electrolyte in the cellreactions. This would allow for a minimum electro-lyte content in the cell, minimizing its weight. Jungnerexperimented with many metals and metal oxides aselectrode materials, including zinc, cadmium, iron,copper oxide, silver oxide, and manganese oxide.

In parallel with Jungner’s efforts in Sweden,Thomas Edison in the United States worked onsimilar ideas using alkaline electrolytes and many ofthe same electrode materials. Patents to these twoinventors were issued at nearly the same time in 1901.

During the period since the beginning of the 20thcentury, a wide variety of cells have been investigated,with many of them being developed into commercial

zz

z

c

cc

FIGURE 2 A set of three Daniell cells connected in series.

Batteries, Overview 119

products for a wide range of applications. Represen-tative cells are discussed in the next section.

3. TYPES OF BATTERIES ANDTHEIR CHARACTERISTICS

Batteries are usually categorized by their ability to berecharged or not, by the type of electrolyte used, andby the electrode type. Examples of various types ofelectrolytes used in batteries are presented in Table I.

Primary (nonrechargeable) cells with aqueouselectrolytes are usually the least expensive and havereasonably long storage lives. Historically, the mostcommon primary cell is the zinc/manganese dioxidecell used for flashlights and portable electronics in sizesfrom a fraction of an amp-hour to a few amp-hours.Various modifications of this cell have been intro-duced, and now the most common version uses analkaline (potassium hydroxide) electrolyte. The mate-rials in this cell are relatively benign from an environ-mental point of view. The electrode reactions in analkaline electrolyte may be represented as follows:


and

2MnO2 þ 2H2O þ 2e� ¼ 2MnOOH þ 2OH�:

The cell potential is 1.5 V, and the theoretical specificenergy is 312 Wh/kg. Practical cells commonly yield40 to 50 Wh/kg.

Another common primary cell is the zinc/air cell.It uses an aqueous potassium hydroxide electrolyte, azinc negative electrode, and a porous catalyzedcarbon electrode that reduces oxygen from the air.The overall electrode reactions are as follows:


and

12O2 þ H2O þ 2e� ¼ 2OH�:

The sum of these reactions gives this overall cellreaction:

Zn þ 12O2 ¼ ZnO:

The cell voltage is 1.6 V, and the theoretical specificenergy is 1200 Wh/kg. Practical specific energy valuesof 200 Wh/kg have been achieved for zinc/air cells.This cell has a very high energy per unit volume,225 Wh/L, which is important in many applications.The materials are inexpensive, so this cell is verycompetitive economically. This cell is quite acceptablefrom an environmental point of view.

The zinc/mercuric oxide cell has the uniquecharacteristic of a very stable and constant voltageof 1.35 V. In applications where this is important,this is the cell of choice. It uses an aqueous potassiumhydroxide electrolyte and is usually used in smallsizes. The electrode reactions are as follows:


and

HgO þ H2O þ 2e� ¼ Hg þ 2OH�:

The high equivalent weight of the mercury results inthe low theoretical specific energy of 258 Wh/kg.Because mercury is toxic, there are significantenvironmental concerns with disposal or recycling.

The zinc/silver oxide cell has a high specific energyof about 100 Wh/kg (theoretical specific ener-gy¼ 430 Wh/kg), and because of this, the cost ofthe silver is tolerated in applications that are notcost-sensitive. Potassium hydroxide is the electrolyteused here. The electrode reactions are as follows:


and

AgO þ H2O þ 2e� ¼ Ag þ 2OH�:

This gives the following overall cell reaction:

Zn þ AgO ¼ Ag þ ZnO:

The cell voltage ranges from 1.8 to 1.6 V duringdischarge. Because of the high density of the silveroxide electrode, this cell is quite compact, givingabout 600 Wh/L.

Primary (nonrechargeable) cells with nonaqueouselectrolytes have been under development since the1960s, following the reports of Tobias and Harrison the use of propylene carbonate as an organicsolvent for electrolytes to be used with alkali metalelectrodes. It has long been recognized that lithiumhas very low equivalent weight and electronegativity(Table II). These properties make it very attractive

TABLE I

Electrolyte Types and Examples

Aqueous electrolyte H2SO4, KOH

Nonaqueous electrolyte

Organic solvent Li salt in ethylene carbonate-diethyl carbonateSolid

Crystalline Na2O�11Al2O3

Polymeric Li salt in polyethylene

Molten salt (high

temperature) LiCl–KCl


for use as the negative electrode in a cell. Becauselithium will react rapidly with water, it is necessaryto use a completely anhydrous electrolyte. There aremany organic solvents that can be prepared free ofwater. Unfortunately, organic solvents generally havelow dielectric constants, making them poor solventsfor the salts necessary to provide electrolyticconduction. In addition, organic solvents are ther-modynamically unstable in contact with lithium, sosolvents that react very slowly and form thin,conductive, protective films on lithium are selected.

Perhaps the most common lithium primary cell isthat of Li/MnO2. It has the advantage of high voltage(compared with aqueous electrolyte cells), 3.05 V, anda high specific energy (B170 Wh/kg). The electro-de reactions are as follows:

Li ¼ Liþ þ e�

and

MnO2 þ Liþ þ e� ¼ LiMnO2:

The electrolyte is usually LiClO4 dissolved in amixture of propylene carbonate and 1,2-dimethoxy-ethane. In general, the specific power that can bedelivered by these cells is less than that available fromaqueous electrolyte cells because the ionic conductiv-ity of the organic electrolyte is much lower than thatof aqueous electrolytes.

An interesting cell is the one using fluorinatedcarbon as the positive electrode material. Because thismaterial is poorly conducting, carbon and titaniumcurrent collection systems are used in the positiveelectrode. The electrode reaction is as follows:

ðCFÞn þ nLiþ þ ne� ¼ nLiF þ nC:

The electrolyte commonly is LiBF4 dissolved ingamma butyrolactone.

An unusual primary cell is that of lithium/thionylchloride (SOCl2). The thionyl chloride is a liquid andcan act as the solvent for the electrolyte salt (LiAlCl4)and as the reactant at the positive electrode. This cellcan function only because of the relatively stable thin

protective film that forms on the lithium electrode,protecting it from rapid spontaneous reaction withthe thionyl chloride. The cell reaction mechanismthat is consistent with the observed products ofreaction is the following:

4Li þ 2SOCl2 ¼ 4LiCl þ 2SO0

and

2SO0 ¼ SO2 þ S;

where SO0 is a radical intermediate that producesSO2 and sulfur. This cell provides a high cell voltageof 3.6 V and a very high specific energy but can beunsafe under certain conditions. Specific energies ofup to 700 Wh/kg have been achieved compared withthe theoretical value of 1460 Wh/kg.

Table III summarizes the characteristics of severalrepresentative primary cells.

Primary cells with molten salt electrolytes arecommonly used in military applications that requirea burst of power for a short time (a few seconds to afew minutes). These batteries are built with anintegral heating mechanism relying on a chemicalreaction to provide the necessary heat to melt theelectrolyte (at 400–5001C). A typical cell uses lithiumas the negative electrode and iron disulfide as thepositive electrode. The following are example elec-trode reactions:

Li ¼ Liþ þ e�

and

FeS2 þ 4Li ¼ 2Li2S þ Fe:

The molten salt electrolyte (a mixture of alkali metalchlorides) has a very high conductivity, allowing thecell to operate at very high specific power levels inexcess of 1 kW/kg.

Rechargeable cells with aqueous electrolytes havebeen available for more than 140 years, beginningwith the Plante cell using lead and sulfuric acid asdiscussed previously. Modern Pb/PbO2 cells andbatteries have received the benefit of many incre-mental improvements in the design and optimizationof the system. Current versions are very reliable andinexpensive compared with competitors. Dependingon the design of the cell, lifetimes vary from a fewyears to more than 30 years. Specific energy values ofup to approximately 40 Wh/kg are available.

Alkaline electrolyte systems are available with avariety of electrodes. Perhaps the most commonalkaline electrolyte cell is the Cd/NiOOH cell. Itoffers very long cycle life (up to thousands of charge–discharge cycles) and good specific power (hundreds

TABLE II

Some Properties of Lithium and Zinc

Lithium Zinc

Equivalent weight (g/equiv.) 6.94 32.69

Reversible potential (V) �3.045 �0.763

Electronegativity 0.98 1.65

Density (g/cm3) 0.53 7.14


of W/kg), albeit with a modest specific energy (35–55 Wh/kg). The cell voltage, 1.2 V, is lower than mostand can be somewhat of a disadvantage. Theelectrode reactions are as follows:

Cd þ 2OH� ¼ CdðOHÞ2 þ 2e�

and

2NiOOH þ 2H2O þ 2e� ¼ 2NiðOHÞ2 þ 2OH�:

Of course, the reverse of these reactions takes placeon recharge.

A very robust rechargeable cell is the Edison cell,which uses iron as the negative electrode, nickeloxyhydroxide as the positive electrode, and anaqueous solution of 30 w/o potassium hydroxide asthe electrolyte. During the early years of the 20thcentury, Edison batteries were used in electric vehicles.They proved themselves to be very rugged anddurable, although they did not have high perfor-mance. The iron electrode reaction is as follows:

Fe þ 2OH� ¼ FeðOHÞ2 þ 2e�:

The positive electrode reaction is the same as for theCd/NiOOH cell described earlier.

The NiOOH electrode has proven itself to begenerally the best positive electrode for use inalkaline electrolytes and has been paired with manydifferent negative electrodes, including Cd, Fe, H2,MH, and Zn. Recently, the metal hydride/nickeloxyhydroxide cell (MH/NiOOH) has been a sig-nificant commercial success and has captured a largefraction of the market for rechargeable cells. It offersa sealed, maintenance-free system with no hazardousmaterials. The performance of the MH/NiOOH cellis very good, providing up to several hundred watts/kilogram peak power, up to about 85 Wh/kg, andseveral hundred cycles. The negative electrodeoperates according to the following reaction:

MH þ OH� ¼ M þ H2O þ e�:

The very high theoretical specific energy and lowmaterials cost of the zinc/air cell make it attractivefor consumer use. There have been many attempts todevelop a rechargeable zinc/air cell, but with limitedsuccess due to the difficulties of producing a high-performance rechargeable air electrode.

The electrode reactions during discharge for thiscell are as follows:


and

O2 þ 2H2O þ 4e� ¼ 4OH�:

During recent years, the cycle life of the airelectrode has been improved to the point where morethan 300 cycles are now feasible. Developmentefforts continue. In the meantime, various versionsof a ‘‘mechanically rechargeable’’ zinc/air cell havebeen tested. These cells have provision for removingthe discharged zinc and replacing it with fresh zinc.This approach avoids the difficulties of operating theair electrode in the recharge mode but creates theneed for recycling the spent zinc to produce new zincelectrode material.

The status of some rechargeable cells withaqueous electrolytes is shown in Table IV.

Rechargeable cells with nonaqueous electrolyteshave been under development for many years,although they have been available on the consumermarket for only a decade or so. All of the types ofnonaqueous electrolytes shown in Table I have beenused in a variety of systems.

Organic solvent-based electrolytes are the mostcommon of the nonaqueous electrolytes and arefound in most of the rechargeable lithium cellsavailable today. A typical electrolyte consists of amixture of ethylene carbonate and ethyl-methylcarbonate, with a lithium salt such as LiPF6 dissolvedin it. Various other solvents, including propylenecarbonate, dimethyl carbonate, and diethyl carbo-nate, have been used. Other lithium salts that havebeen used include lithium perchlorate, lithiumhexafluoro arsenate, and lithium tetrafluoro borate.All of these combinations of solvents and salts yieldelectrolytes that have much lower conductivities thando typical aqueous electrolytes. As a result, theelectrodes and electrode spacing in these lithium cellsare made very thin to minimize the cell resistance andmaximize the power capability.

Lithium is difficult to deposit as a smoothcompact layer in these organic electrolytes, so a hostmaterial, typically carbon, is provided to take up the

TABLE III

Properties of Some Lithium Primary Cells

Open circuit

voltage

(V)

Practical specific

energy

(Wh/kg)

Theoretical

specific energy

(Wh/kg)

Li/MnO2 3.0 280 750

Li/(CF)n 2.8 290 2100

Li/SOCl2 3.6 450–700 1460

Li/FeS2 1.7 220 760


Li on recharge and deliver it as Li ions to theelectrolyte during discharge. Many types of carbon,both graphitic and nongraphitic, have been used asthe lithium host material. In addition, a variety ofintermetallic compounds and metals have been usedas Li host materials. All of the commercial Lirechargeable cells today use a carbon host materialas the negative electrode.

The positive electrodes of Li cells are usually metaloxide materials that can intercalate Li with aminimal change in the structure of the material.The most common positive electrode material used incommercial Li cells is LiCoO2. This material is verystable toward the removal and reinsertion of Li intoits structure. It can be cycled more than 1000 timeswith little change in capacity.

The electrode reactions during discharge for thiscell are as follows:

xLiðCÞ ¼ xLiþ þ C þ xe�

and

Li1�xCoO2 þ xLiþ þ xe� ¼ LiCoO2:

Excellent engineering of this cell has resulted in aspecific energy of more than 150 Wh/kg and a peakspecific power of more than 200 W/kg, making itvery desirable for applications requiring high specificenergy. This cell has been manufactured extensivelyfirst in Japan and now in Korea and China, and it hasbeen used worldwide.

Even though the Li(C)/CoO2 cell has manydesirable features, it has some limitations anddisadvantages. The use of Co represents an environ-mental hazard because Co is toxic. It is also tooexpensive to use in large batteries suitable for low-cost applications such as electric and hybrid vehiclepropulsion. In addition, there are safety problems

associated with overcharge of this cell. Significantovercharge can cause the cell to vent and burn.Operation at elevated temperatures can also causefire. Special microcircuits mounted on each cellprovide control over the cell, preventing overchargeand overdischarge.

During the past several years, there has been alarge effort to replace the LiCoO2 with a lessexpensive, more environmentally benign materialthat can withstand overcharge and overdischargewhile retaining the excellent performance of LiCoO2.Many metal oxide materials have been synthesizedand evaluated. Some of them are promising, but nonehas surpassed LiCoO2 so far. Popular candidatematerials include lithium manganese oxides such asLiMn2O4, LiMnO2, and materials based on them,with some of the Mn replaced by other metals suchas Cr, Al, Zn, Ni, Co, and Li. Other candidatesinclude LiNiO2, LiFePO4, and variants on these,with some of the transition metal atoms replaced byother metal atoms in an attempt to make thematerials more stable to the insertion and removalof the Li during discharge and charge. Progress isbeing made, and some Mn-based electrode materialsare appearing in commercial cells.

Some of the instabilities that have been observedare traceable to the electrolyte and its reactivity withthe electrode materials. This has resulted in someresearch on more stable solvents and Li salts. The Lisalts must be stable to voltages approaching 5 V andmust be very soluble and highly conductive. Thesolvents must have a wide temperature range ofliquidity so that cells can be operated from very coldto very warm temperatures (�40 to þ 601C for someapplications).

The emphasis on lower cost electrode materialshas continued, and some encouraging results havebeen obtained for materials based on LiFePO4. This

TABLE IV

Rechargeable Aqueous Battery Status

System Cell voltage (V)

Theoretical specific

energy (Wh/kg)

Specific energy

(Wh/kg)

Specific power

(W/kg) Cycle life

Cost

(dollars/kWh)

Pb/PbO2 2.10 175 30–45 50–100 4700 60–125

Cd/NiOOH 1.20 209 35–55 400 2000 4300

Fe/NiOOH 1.30 267 40–62 70–150 500–2000 4100

H2/NiOOH 1.30 380 60 160 1000–2000 4400

MH/NiOOH 1.20 200 60–85 200þ 500–600 4400

Zn/NiOOH 1.74 326 55–80 200–300 500þ 150–250 (est.)

Zn/Air 1.60 1200 65–120 o100 300 100 (est.)

Note. est., estimate.


material is very stable toward repeated insertion andremoval of Li and has shown hundreds of cycles withlittle capacity loss. One problem with LiFePO4 is itslow electronic conductivity, making full use and highrate operation difficult, even with enhanced currentcollection provided by added carbon.

Some unusual electrode materials are underinvestigation in the quest to reduce cost. Theseinclude elemental sulfur and organo-disulfide poly-mers. The Li/S cell has been investigated severaltimes during the past 20 years or more. From atheoretical specific energy point of view, it is veryattractive at 2600 Wh/kg. The cost of sulfur isextremely low, and it is environmentally benign. Asignificant difficulty with the sulfur electrode in mostorganic solvents is the formation of soluble poly-sulfides during discharge. The soluble polysulfidescan migrate away from the positive electrode andbecome inaccessible, resulting in a capacity loss.

The overall reaction for the sulfur electrode is asfollows:

S þ 2e� ¼ S�2:

And the overall cell reaction is as follows:

2Li þ S ¼ Li2S:

Recent work on this system has involved isolatingthe soluble polysulfides with an ionically conductivemembrane or using solvents that have a lowsolubility for the polysulfides that are formed duringdischarge. These approaches have resulted in a largeincrease of cycle life, up to several hundred cycles.

The organo–disulfides offer a lower specific energythan does elemental sulfur, but they do not exhibitthe same relatively high solubility. Long cycle liveshave been demonstrated for these electrodes. They

are usually coupled with a lithium metal negativeelectrode, which has a short cycle life. Current effortsare under way to improve the cycle life of the lithiummetal electrode.

The status of some rechargeable cells withnonaqueous electrolytes is presented in Table V.The performance of several rechargeable batteries isshown in Fig. 3.

There are a few rechargeable batteries withmolten salt or ceramic electrolytes that operate atelevated temperatures, but these have found onlyspecialty (demonstration) applications such as sta-tionary energy storage and utility vehicle propulsionand are not yet in general use. Examples of these are

TABLE V

Nonaqueous Rechargeable Battery Status

System

Cell voltage

(V)

Equivalent

weight

(kg)

Theoretical

specific energy

(Wh/kg)

Specific energy

(Wh/kg)

Specific power

(W/kg)

Cost

(dollars/kWh)

Li/CoO2 3.60 0.1693 570 125 4200 44100

Li/Mn2O4 4.00 0.1808 593 150 200 B200

Li/FePO4 3.5 0.151 621 120 100(est.) B200

Li/V6O13 2.40 0.0723 890 150 200 4200

Li/TiS2 2.15 0.1201 480 125 65 4200

Li/SRS 3.00 0.080 1000 200(est.) 400(est.) N.A.

Li/S 2.10 0.0216 2600 300 200 o100

Note. est., estimate; N.A., not available.

Internalcombustionengine

Li ion cells

Na/S

Zn/Air

Pb/PbO2

MH/NiOOH

Zn/NiOOH

Cd/NiOOH

1000

100

1010 100 1000

Wh/kg

W/k

g

FIGURE 3 Specific power, watts/kilogram (W/kg), as a func-tion of specific energy, watt-hours/kilogram (Wh/kg), for various

cells. A curve for an internal combustion engine is included for

reference.


the lithium/iron sulfide cell with a molten alkalihalide electrolyte, the sodium/sulfur cell with a beta-alumina ceramic (Na2O .Al2O3) electrolyte, and thesodium/nickel chloride cell with both molten salt andbeta-alumina electrolytes. These cells operate at 350to 4001C. The characteristics of these cells are shownin Table VI.

4. APPLICATIONS

Current applications of batteries are very wide-ranging. It is difficult to identify many technologiesthat do not rely on either primary or secondarybatteries. Batteries are part of our everyday lives andare essential to many fields with which most peopledo not come into contact.

Consumer applications of batteries include flash-lights, portable electronics (e.g., cell phones), radios,televisions, MP3 players, personal digital assistants,notebook computers, portable tools, power gardentools, burglar alarms, smoke detectors, emergencylighting, remote control devices for television,stereos, electronic games, remote-controlled toys,and other entertainment equipment. Medical appli-cations include heart pacemakers, medical diagnosticdevices (e.g., glucose monitors), pumps for adminis-tering drugs, and nerve and muscle stimulators.Automobiles depend on batteries for starting, light-ing, ignition, and related functions. Golf carts, utilityvehicles, and forklift trucks represent a variety ofpropulsion-power applications.

There are many government and military applica-tions with stringent requirements for batteries,including communications radios and phones, por-table lasers for range finding, field computers, nightvision equipment, sensors, starting for all types ofvehicles, power sources for fuses, timers, variousmissiles, aircraft starting and electrical systems,electrically powered vehicles (e.g., submarines), andoff-peak energy storage for utility power.

The space program depends critically on batteriesfor all manner of power supplies for communication,operation of the electrical systems in shuttle craft andspace stations, power for the small vehicles andsensors deployed on the moon and planets, life supportsystems, power for space suits, and many others.

The batteries that are used in these applicationscome in many sizes, shapes, and types. The sizes ofindividual cells range from a couple of millimeters orless to a couple of meters. Depending on the appli-cation, many cells are connected in series-parallelarrangements to provide the required voltage andcurrent for the necessary time.

Future applications for batteries are certain toinclude those listed here as well as a variety that aredifficult to imagine at this time. We are now seeingthe appearance of electric hybrid vehicles with abattery as part of the propulsion system and anelectric motor. These vehicles are much more efficientthan the standard vehicles and deliver approximately50 miles per gallon. In this application, metalhydride/nickel oxide cells and lithium cells are used.We are certain to see many more of these hybridvehicles with relatively large batteries providing muchof the propulsion power. These batteries store a fewkilowatt-hours of energy. We have also seen a smallnumber of all-battery electric vehicles on the road.They use larger batteries of up to 20 kWh storagecapability. As batteries of higher specific energy andlower cost become available, more electrically pow-ered vehicles (both battery and hybrid) will be sold.

Systems for the remote generation and storage ofelectrical energy depend on batteries. These includesolar- and wind-powered systems for remote build-ings. We can expect to see ever-smaller cells beingused as power sources on microcircuit chips incomputers and other electronic devices. Biomedicalapplications will expand with a wide variety ofbattery-powered sensors, stimulators, transmitters,and the like. A battery-powered artificial heart is adistinct possibility.

TABLE VI

High-Temperature Battery Status

System

Cell voltage

(V)

Theoretical

specific energy

(Wh/kg)

Specific energy

(Wh/kg)

Specific power

(W/kg) Cycle life

Cost

(dollars/kWh)

Temperature

(1C)

Na/S 2.08 758 180 160 500–2000 4100 350

Na/NiCl2 2.58 787 90 90–155 600 4300 350

LiAl/FeS2 1.80 514 180–200 4200 1000 N.A. 400

Note. N.A., not available.


Future household applications include a battery-powered robot capable of performing routine house-hold tasks. As computer-controlled homes becomemore common, batteries will be used in manysensors, transmitters, and receivers, either becausethey are mobile or to be independent of the powergrid. As fuel cells are introduced, it is likely that inmost applications the fuel cell will be accompaniedby a battery to store energy and meet peak powerrequirements, lowering the total cost of the systemfor a given peak power requirement.

5. ENVIRONMENTAL ISSUES

Some types of batteries contain hazardous materialsthat must be kept out of the environment. Examplesof these are the lead in lead–acid batteries and thecadmium in cadmium/nickel oxide batteries. Asbatteries continue to experience growth in their use,developers and manufacturers must be cognizant ofenvironmental issues relating to the materials used inthe batteries. For example, more than 80% of the leadin a new lead–acid battery has been recycled. With theuse of clean technology in the recycling process, thishazardous material can be used and reused safely.

6. FUTURE OUTLOOK

We can expect to see continuing growth in the varietyof electrochemical couples available and the sizes andshapes of cells to satisfy a growing variety ofapplications, from microwatt-hours to megawatt-hours of energy. The world market for batteries is amulti-billion-dollar per year market and has beengrowing very rapidly, especially during the pastdecade or so. This growth will continue, especially

as the developing countries install their communica-tions systems and more extensive infrastructures forelectrical power and transportation. Lower costbatteries will be in demand, as will batteries forvarious specialty applications. Batteries capable offitting into oddly shaped spaces will be needed, andcells with polymer electrolytes can satisfy this uniquerequirement. Such batteries are being introduced now.

As batteries show an increase in specific energy inthe range of 300 to 400 Wh/kg, battery-poweredelectric cars will become quite attractive. Batteriescould power buses and other vehicles with ranges of200 to 300 miles.

In the development of new cells, it will beimportant to continue to reduce the equivalent weightof the electrodes. At this time, the highest theoreticalspecific energy of any cell being developed is that of Li/S at 2600 Wh/kg. If a lithium/oxygen, lithium/phos-phorus, or lithium/fluorine cell could be developed,this would lead to an even higher specific energy.

SEE ALSO THEFOLLOWING ARTICLES

Batteries, Transportation Applications � ElectricalEnergy and Power � Electric Motors � Storage ofEnergy, Overview

Further Reading

Linden, D., and Reddy, T. B. (eds.). (2002). ‘‘Handbook of

Batteries,’’ 3rd ed. McGraw–Hill, New York.

Socolow, R., and Thomas, V. (1997). The industrial ecology of

lead and electric vehicles. J. Industr. Ecol. 1, 13.Vincent, C. A., Bonino, F., Lazzari, M., and Scrosati, B. (1984).

‘‘Modern Batteries: An Introduction to Electrochemical Power

Sources.’’ Edward Arnold, London.


Documents

Batteries, Overview